Vintage Sacks Read Free

expanded, forming a monstrous pillar of carbon rising high above the rim of the beaker. “Beware,” David said, as I gazed at this transformation. “
You’ll
be turned into a pillar of carbon if you get the acid on yourself.” And then he told me horror stories, probably invented, of vitriol throwings in East London, and patients he had seen coming into the hospital with their entire faces all but burned off. (I was not quite sure whether to believe him, for when I was younger he had told me that if I looked at the Kohanim as they were blessing us in the shul—their heads were covered with a large shawl, a tallis, as they prayed, for they were irradiated, at this moment, by the blinding light of God—my eyes would melt in their sockets and run down my cheeks like fried eggs.) 1
    I spent a good deal of my time in the lab examining chemical colors and playing with them. There were certain colors that held a special, mysterious power for me—this was especially so of very deep and pure blues. As a child I had loved the strong, bright blue of the Fehling’s solution in my father’s dispensary, just as I had loved the cone of pure blue at the center of a candle flame. I found I could produce very intense blues with some cobalt compounds, with cuprammonium compounds, and with complex iron compounds like Prussian blue.
    But the most mysterious and beautiful of all the blues for me was that produced by dissolving alkali metals in liquid ammonia (Uncle Dave showed me this). The fact that metals
could
be dissolved at all was startling at first, but the alkali metals were all soluble in liquid ammonia (some to an astounding degree—cesium would completely dissolve in a third its weight of ammonia). When the solutions became more concentrated, they suddenly changed character, turning into lustrous bronze-colored liquids that floated on the blue—and in this state they conducted electricity as well as liquid metal like mercury. The alkaline earth metals would work as well, and it did not matter whether the solute was sodium or potassium, calcium or barium—the ammoniacal solutions, in every case, were an identical deep blue, suggesting the presence of some substance, some structure, something common to them all. It was like the color of the azurite in the Geological Museum, the very color of heaven.
    Many of the so-called transition elements infused their compounds with characteristic colors—most cobalt and manganese salts were pink; most copper salts deep blue or greenish blue; most iron salts pale green and nickel salts a deeper green. Similarly, in minute amounts, transition elements gave many gems their particular colors. Sapphires, chemically, were basically nothing but corundum, a colorless aluminum oxide, but they could take on every color in the spectrum—with a little bit of chromium replacing some of the aluminum, they would turn ruby red; with a little titanium, a deep blue; with ferrous iron, green; with ferric iron, yellow. And with a little vanadium, the corundum began to resemble alexandrite, alternating magically between red and green—red in incandescent light, green in daylight. With certain elements, at least, the merest smattering of atoms could produce a characteristic color. No chemist could have “flavored” corundum with such delicacy, a few atoms of this, a few ions of that, to produce an entire spectrum of colors.
    There were only a handful of these “coloring” elements—titanium, vanadium, chromium, manganese, iron, cobalt, nickel, and copper, so far as I could see, being the main ones. They were, I could not help noticing, all bunched together in terms of atomic weight—though whether this meant anything, or was just a coincidence, I had no idea at the time. It was characteristic of all of these, I learned, that they had a number of possible valency states, unlike most of the other elements, which had only one. Sodium, for